Atomic physics

Electronic configuration

Electrons form notional shells around the nucleus. These are normally in a ground state but can be excited by the absorption of energy from light (photons), magnetic fields, or interaction with a colliding particle (typically ions or other electrons).

Electrons that populate a shell are said to be in a bound state. The energy necessary to remove an electron from its shell (taking it to infinity) is called the binding energy. Any quantity of energy absorbed by the electron in excess of this amount is converted to kinetic energy according to the conservation of energy. The atom is said to have undergone the process of ionization.

If the electron absorbs a quantity of energy less than the binding energy, it will be transferred to an excited state. After a certain time, the electron in an excited state will "jump" (undergo a transition) to a lower state. In a neutral atom, the system will emit a photon of the difference in energy, since energy is conserved.

If an inner electron has absorbed more than the binding energy (so that the atom ionizes), then a more outer electron may undergo a transition to fill the inner orbital. In this case, a visible photon or a characteristic X-ray is emitted, or a phenomenon known as the Auger effect may take place, where the released energy is transferred to another bound electron, causing it to go into the continuum. The Auger effect allows one to multiply ionize an atom with a single photon.

There are rather strict selection rules as to the electronic configurations that can be reached by excitation by light –however, there are no such rules for excitation by collision processes.

Bohr model of the atom

The Bohr model, proposed by Niels Bohr in 1913, is a revolutionary theory describing the structure of the hydrogen atom. It introduced the idea of quantized orbits for electrons, combining classical and quantum physics.

Key Postulates of the Bohr Model

1. Electrons Move in Circular Orbits
   • Electrons revolve around the nucleus in fixed, circular paths called orbits or energy levels.
   • These orbits are stable and do not radiate energy.
2. Quantization of Angular Momentum:
   • The angular momentum of an electron is quantized and given by: $${\displaystyle L=m_{\text{e}}vr=n\hslash ,n=1,2,3,\dots }$$ where:

     ${\displaystyle m_{\text{e}}}$: electron mass

     ${\displaystyle v}$: velocity of the electron

     ${\displaystyle r}$: radius of the orbit

     ${\displaystyle \hslash }$: reduced Planck constant (${\displaystyle \hslash =h/2\pi }$)

     ${\displaystyle n}$: principal quantum number, representing the orbit

3. Energy Levels
   • Each orbit has a specific energy. The total energy of an electron in the ${\displaystyle n}$th orbit is: $${\displaystyle E_n=-\frac{13.6\mathrm{e}\mathrm{V}}{n^2},}$$ where ${\displaystyle 13.6\mathrm{e}\mathrm{V}}$ is the ground-state energy of the hydrogen atom.
4. Emission or Absorption of Energy
   • Electrons can transition between orbits by absorbing or emitting energy equal to the difference between the energy levels: $${\displaystyle \mathrm{\Delta }E=E_{\text{f}}-E_{\text{i}}=h\nu ,}$$ where:

     ${\displaystyle h}$: the Planck constant.

     ${\displaystyle \nu }$: frequency of emitted/absorbed radiation.

     ${\displaystyle E_{\text{f}},E_{\text{i}}}$: final and initial energy levels.

History and developments

One of the earliest steps towards atomic physics was the recognition that matter was composed of atoms. It forms a part of the texts written in 6th century BC to 2nd century BC, such as those of Democritus or Vaiśeṣika Sūtra written by Kaṇāda. This theory was later developed in the modern sense of the basic unit of a chemical element by the British chemist and physicist John Dalton in the 18th century. At this stage, it was not clear what atoms were, although they could be described and classified by their properties (in bulk). The invention of the periodic system of elements by Dmitri Mendeleev was another great step forward.

The true beginning of atomic physics is marked by the discovery of spectral lines and attempts to describe the phenomenon, most notably by Joseph von Fraunhofer. The study of these lines led to the Bohr atom model and to the birth of quantum mechanics. In seeking to explain atomic spectra, an entirely new mathematical model of matter was revealed. As far as atoms and their electron shells were concerned, not only did this yield a better overall description, i.e. the atomic orbital model, but it also provided a new theoretical basis for chemistry (quantum chemistry) and spectroscopy.

Since the Second World War, both theoretical and experimental fields have advanced at a rapid pace. This can be attributed to progress in computing technology, which has allowed larger and more sophisticated models of atomic structure and associated collision processes. Similar technological advances in accelerators, detectors, magnetic field generation and lasers have greatly assisted experimental work.

Beyond the well-known phenomena which can be described with regular quantum mechanics chaotic processes can occur which need different descriptions.

Significant atomic physicists

Pre quantum mechanics

• John Dalton
• Joseph von Fraunhofer
• Johannes Rydberg
• J. J. Thomson
• Ernest Rutherford
• Democritus
• Vaiśeṣika Sūtra

Post quantum mechanics

• Alexander Dalgarno
• David Bates
• Niels Bohr
• Max Born
• Clinton Joseph Davisson
• Paul A. M. Dirac
• Enrico Fermi
• Charlotte Froese Fischer
• Vladimir Fock
• Douglas Hartree
• Ernest M. Henley
• Ratko Janev
• Daniel Kleppner
• Harrie S. Massey
• Nevill Mott
• I. I. Rabi
• Norman Ramsey
• Mike Seaton
• John C. Slater
• George Paget Thomson